Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. Assemble models A–K from instruction sheet.
2. Make copies of the periodic table and Data Tables A and B for each student.

Safety Precautions

Although this activity is considered nonhazardous, observe all normal laboratory safety guidelines.

Lab Hints

• This is a paper-and-pencil activity—two class sessions are recommended for its completion. Students may work on Lewis structures (Part A) one day, examine and sketch models (Part B) in class another day, and then complete the sketches of the molecules needed for Part A out of class, if desired. Since this is a dry-lab exercise, students may work individually. Having students work collaboratively in groups of two or three, however, will provide a better learning environment, making it easier for students to brainstorm and run ideas back and forth.
• The molecular models required for Part B are pictured in the Supplemental Information of the Teacher PDF. Many different kinds of model sets may be used to build additional molecular models required for Part B. Inexpensive homemade models may also be constructed using various size Styrofoam balls and pipe cleaners or wires. See the Supplementary Information for photos of the models required for this activity. If enough model sets are available, consider having students build their own models.
• Additional plastic models can be built using the Shapes of Molecules Model Set (Catalog No. AP5456) and the Inorganic/Organic Teacher Model Set (Catalog No. AP5455) available from Flinn Scientific. Replacement atoms and bonds for the Inorganic/Organic Model Set are also available. The “Shapes of Molecules” set may be used to prepare one model for each of the principal types of molecular geometries (linear, bent, trigonal planar, pyramidal, tetrahedral, trigonal bipyramidal, octahedral, square pyramidal and square planar). With the Inorganic/Organic Teacher Model Set, it is possible to represent double and triple bonds as distinct from single bonds.
• Because copies of Data Table A and B use up a lot of paper, teachers may prefer to have students prepare their own tables. The sample tables may be used as templates—the teacher may wish to prepare overhead transparencies of Data Tables A and B.

Teacher Tips

• According to the most recent, official IUPAC recommendations, the columns in the periodic table are numbered continuously from 1–18, with no breaks for the transition metals. The A/B Roman numeral numbering system (see Figure 1 in the Background section) is often still shown in textbooks, probably because it helps students predict valence electrons and ionic charges.

• Many teachers have developed their own protocols for drawing Lewis structures. There are no hard-and-fast rules. One approach that works well is to study the typical number of bonds that an atom will form when it has zero formal charge (e.g., carbon forms four bonds, nitrogen three bonds, oxygen two bonds, fluorine one bond). Drawing Lewis structures always requires trial-and-error, however, no matter how detailed the “rules” that the students are given.
• Students may struggle with how the VSEPR model is used to predict the shapes of molecules containing multiple bonds (see Part B, step 4). Each multiple bond is treated as though it were a single electron pair. This is because all electron pairs of a multiple bond are required to be in approximately the same region of space.
• Explain to students that lone pair electrons, as a rule, require more space than bonded pairs. When looking at molecules that can have more than one spatial arrangement of atoms (i.e., XeF₄, select the geometry that allows the lone pairs the greatest separation).
• Having students sketch three-dimensional structures of molecules is an important part of this activity. The spatial reasoning skills required to visualize molecules in three dimensions will vary greatly across the student population. All students will benefit from the opportunity to hold models in their hands, rotate them, turn them upside down, etc. Some students, for example, will have no trouble “following the rules” and drawing Lewis structures for molecules. They may need more time, however, to see the molecules in three dimensions and to identify why different molecules share the same molecular geometry. Another group of students may proceed more slowly through the Lewis structures but will be able to recognize three-dimensional structures with ease.
• Students may refer to their textbooks for additional examples, drawings, and models of molecular geometry. We recommend, however, that students start with the actual physical models used in Part B to learn about VSEPR theory and molecular geometry. Words and pictures on the printed page are not an adequate substitute for students holding the models in their hands as they try to visualize the shapes of molecules.

Answers to Prelab Questions

1. Write the Lewis electron-dot symbol for each of the following atoms: hydrogen, boron, nitrogen, silicon, sulfur and bromine.
   {12720_Answers_Figure_11}

   Note: There are no rules for where the dots are drawn.

2. What information about a molecule does its Lewis structure provide? What information is neither shown nor implied in the Lewis structure?

   The Lewis structure shows all of the atoms in a molecule and how they are connected via single, double or triple bonds. It also shows any unshared pairs of valence electrons on each atom in the structure. The Lewis structure does not provide any information concerning the three-dimensional structure of the molecule or bond angles between atoms. The structures are drawn in two dimensions and are not meant to be perspective drawings.

3. There are several exceptions to the octet rule.
   1. Based on its electron configuration, explain why hydrogen can only have two valence electrons around it when it bonds to other atoms. What is the maximum number of bonds hydrogen will form?

      Hydrogen has one valence electron in a 1s orbital. The 1s orbital can accommodate only two electrons and there are no p orbitals in the n = 1 principal energy level. Therefore, hydrogen can have only two electrons around it when it bonds to other atoms. Hydrogen can form a maximum of one (single) covalent bond.

   2. Neutral compounds of boron may be described as “electron-deficient.” Based on its electron configuration, predict how many covalent bonds boron will form. Is this the maximum number of bonds boron will form? Hint: Boron forms polyatomic ions.

      The valence electron configuration of boron is 2s²2p¹. Boron has three valence electrons available for bonding and thus should form only three covalent bonds by sharing these electrons with other atoms. In neutral compounds, if one “counts” bonding electrons as belonging to both atoms in the bond, then boron would have six electrons around it, not eight. Because the outermost n = 2 principal energy level for boron will accommodate a total of eight electrons, it is possible for boron to form four covalent bonds if it “accepts” additional electrons. The borohydride (BH₄–) and fluoroborate (BF₄–) anions are examples of polyatomic ions in which boron forms four bonds.

   3. Many elements in the third row and beyond in the periodic table may form more than four bonds and thus appear to have "expanded octets.” Phosphorus and sulfur, for example, may form five and six covalent bonds, respectively. Count up the total number of valence electrons in PCl₅ and draw its Lewis structure. How many valence electrons are "counted" toward the central P atom?

      PCl₅ has a total of 40 valence electrons distributed as shown in the Lewis structure at the right. The phosphorus atom forms five bonds to chlorine atoms and by the electron counting scheme appears to have 10 electrons in its valence shell. This is permitted because the P atom has empty 3d orbitals.

Sample Data

Data Table A

Data Table A (continued)

Data Table A (continued)

Data Table B

Teacher Handouts

12720_Teacher1.pdf

References

We are grateful to Roxie Allen, St. John’s School, Houston TX, for providing Flinn Scientific with the idea and instructions for this activity.

This activity was adapted from Flinn ChemTopic™ Labs, Volume 5, Chemical Bonding, Cesa, I., Ed., Flinn Scientific, Batavia, IL, 2004.

Introduction

Molecules have shape! The structure and shape of a molecule influences its physical properties and affects its chemical behavior as well. Lewis structures and VSEPR theory offer useful models for visualizing the structures of covalent compounds.

Concepts

• Valence electrons

• Covalent bonding
• Lewis structures
• VSEPR theory

Background

Covalent bonds are defined as the net attractive forces between nonmetal atoms that share one, two or three pairs of electrons. In general, only the valence electrons, those in the highest energy levels that are farthest away from the nucleus, are available for bonding. The number of valence electrons influences the number of bonds that an atom will form. The periodic table offers a convenient shortcut for determining the number of valence electrons in an atom. Remember that the position of an element in the modern periodic table reflects its electron configuration. When the representative elements are arranged in columns from IA to VIIIA (see Figure 1), the number of valence electrons for an element is equal to its group number. Thus, potassium in Group IA has one valence electron, carbon in Group IVA has four valence electrons, and chlorine in Group VIIA has seven valence electrons.

In 1916, G. N. Lewis, an American chemist, proposed arranging dots around the symbols of the elements to represent the valence electrons. Lewis electron-dot symbols (see Figure 2) remain the most popular way to illustrate the valence electrons that are available for bonding. Lewis structures build on the Lewis electron-dot symbols of the elements to show the bonding arrangement of atoms in a molecule and the distribution of all valence electrons. The Lewis structure of a molecule thus shows all of the atoms and how they are connected. A single covalent bond between two atoms, corresponding to a pair of electrons, is represented using a dash (—). Sometimes atoms share more than one pair of electrons between them in order to form stable molecules. Two dashes, corresponding to two pairs of electrons and three dashes, corresponding to three pairs of electrons, are used to represent double and triple bonds, respectively.

G. N. Lewis offered a simple theory, based on the known stability of the noble gases (e.g., He, Ne), to predict how many bonds an atom will form and how many atoms of a particular type will come together to form a stable molecule. According to Lewis, nonmetals may share electrons in order to achieve a valence shell electron “count” similar to that of the noble gases:

“Two atoms may conform to the rule of eight, or the octet rule… by sharing one or more pairs of electrons. The electrons which are held in common by two atoms may be considered to belong to the outer shell of both atoms.”

The noble gases have filled s and p orbitals with eight electrons. The octet rule assumes that atoms form molecules to achieve this stable, noble gas electron configuration. In counting valence electrons to predict the structure of a covalent compound, we will distinguish between two kinds of electron pairs. Bonding pairs of electrons are shared between atoms and thus “belong“ to both atoms in the bond. Nonbonding or unshared pairs of electrons are not shared between atoms and are therefore “counted“ toward only one of the atoms. Consider the fluorine molecule (F₂). Each fluorine atom has seven valence electrons and needs only one more electron to form a stable molecule—two fluorine atoms come together and share one bonding pair of electrons (Equation 1).
Each fluorine atom retains its three unshared pairs of electrons. Molecular Geometry

According to the Valence Shell Electron Pair Repulsion (VSEPR) theory, the valence electron pairs that surround an atom repel each other due to their like negative charges. In order to minimize this repulsion, the electron pairs should be positioned around the atom so that they are as far apart as possible. The resulting symmetrical arrangement of electron pairs around atoms can be used to predict molecular geometry—the three-dimensional shape of a molecule. Two pairs of electrons around an atom should adopt a linear arrangement, three pairs a trigonal planar arrangement and so on.

The three-dimensional structure of a molecule is affected by the spatial arrangement of all the electron pairs—both bonding and nonbonding—around the central atom. However, only the physical arrangement of the atoms is used to describe the resulting molecular geometry. This is best illustrated using an example. The Lewis structure of the water molecule is shown as the first example in Figure 3—there are four pairs of valence electrons around the central oxygen atom. Two pairs of electrons are involved in bonding to hydrogen atoms, while the other two electron pairs are unshared pairs. Four pairs of electrons around an atom will adopt a tetrahedral arrangement in space, as depicted in the second example in Figure 3, to be as far apart in space as possible. For this representation, the symbol shows one lone pair of electrons extending behind the plane of the paper. The symbol shows one lone pair of electrons extending in front of the plane of the paper while the symbols “——” represent the hydrogen–oxygen bonds positioned in the plane of the paper. As a result, the two hydrogen atoms and the oxygen atom occupy a “bent” (inverted-V) arrangement.

When two atoms are linked via a double or triple bond (with two or three bonding pairs of electrons, respectively), the multiple electron pairs between the atoms must be considered together when determining the shape of the molecule. Carbon dioxide provides a good example (see Figure 4). The central carbon atom is linked to two oxygen atoms by two double bonds. The resulting arrangement of atoms is linear—both electron pairs in each double bond are considered to be one electron group that must be in approximately the same region, near the oxygen atom.

Experiment Overview

The purpose of this activity is to practice drawing Lewis structures of molecules and to use these structures to predict their molecular geometry. Molecular models will be studied to visualize the shapes of molecules and to sketch their three-dimensional structures.

Materials

Prelab Questions

1. Write the Lewis electron-dot symbol for each of the following atoms: hydrogen, boron, nitrogen, silicon, sulfur and bromine.
2. What information about a molecule does its Lewis structure provide? What information is neither shown nor implied in the Lewis structure?
3. There are several exceptions to the octet rule.
   1. Based on its electron configuration, explain why hydrogen can only have two valence electrons around it when it bonds to other atoms. What is the maximum number of bonds hydrogen will form? 
   2. Neutral compounds of boron may be described as “electron-deficient.” Based on its electron configuration, predict how many covalent bonds boron will form. Is this the maximum number of bonds boron will form? Hint: Boron forms polyatomic ions. 
   3. Many elements in the third row and beyond in the periodic table may form more than four bonds and thus appear to have “expanded octets.” Phosphorus and sulfur, for example, may form five and six covalent bonds, respectively. Count up the total number of valence electrons in PCl₅ and draw its Lewis structure. How many valence electrons are “counted” toward the central P atom?

Safety Precautions

Although this activity is considered non-hazardous, observe all normal laboratory safety guidelines.

Procedure

Part A. Lewis Structures

1. Write the formula for each molecule or polyatomic ion listed in Data Table A. Some of the formulas have been filled in for you.
2. Count the number of valence electrons supplied by each atom in the formula. Determine the total number of valence electrons and record this number in Data Table A. In the case of polyatomic ions, add one electron for each unit of negative charge or subtract one electron for each unit of positive charge to determine the total number of valence electrons.
3. Use the following guidelines to draw a reasonable Lewis structure for each molecule or ion in the space provided. If more than one Lewis structure is reasonable, draw all of the appropriate Lewis structures. Note: Keep in mind the exceptions to the octet rule discussed in the Prelab Questions.
   • Draw a “skeleton” structure for the molecule or ion, joining atoms by single bonds (see Figure 8). If there is a single atom of one element in a compound, show it as the central atom with other atoms joined to it. The least electronegative atom is usually the central atom. However, hydrogen is never a central atom.
     {12720_Procedure_Figure_8}
   • From the total number of valence electrons, subtract two for each single bond in the skeleton—this tells you how many valence electrons are left to distribute.
   • Use the octet rule to distribute the remaining valence electrons as unshared pairs around the atoms in the molecule or ion.
   • If this point is reached and there are too few valence electrons to give each atom an octet, multiple bond(s) may be needed. Remember that bonding electrons are “counted” toward both atoms in the bond while unshared electrons are “assigned” to only one atom.
   • If all else fails, some Lewis structures can only be drawn by assuming there is an unpaired electron in the molecule.

Part B. Molecular Models

4. Examine the molecular models A–K. For each model, identify the number of bonding pairs, the number of unshared pairs, and the total number of electron pairs around the central atom. Record this information in Data Table B. Note: In the case of double or triple bonds, count all of the electrons involved in the bond as one pair of electrons. (Review the structure of carbon dioxide in the Background section.)
5. Sketch the three-dimensional arrangement of valence electron pairs around each central atom in Data Table B. Recall that multiple pairs of bonding electrons in double and triple bonds must "point" to the same atom.
6. Use the following terms to describe the molecular geometry for each model: Linear, trigonal planar, bent, tetrahedral, pyramidal, trigonal bipyramidal, octahedral, square pyramidal and square planar. Record the molecular geometry in Data Table B. Hint: Molecular geometry describes the physical arrangement of the atoms, not the electron pairs.
7. Return to the molecules and polyatomic ions listed in Data Table A. Refer to the models from Part B for comparison: Count the number of valence electron pairs (see step 4) around the central atom and predict the molecular geometry for each molecule or ion in Data Table A.
8. Draw a three-dimensional sketch of the molecule or ion in the space below its name and describe its molecular geometry (see step 6) in Data Table A. Example: The structure of H₂O could be sketched as follows:
   {12720_Procedure_Figure_9}
   The structure of H₃O⁺ could be sketches as follows:
   {12720_Procedure_Figure_10}

Student Worksheet PDF

12720_Student1.pdf